General Chemistry II: Exam II Study Guide and Practice Questions

Chapter 3: Chemical Bonding and Molecular Structure

Ionic Bonding

Ionic bonding involves the electrostatic attraction between oppositely charged ions, typically formed between metals and nonmetals.

• Coulomb Stabilization Energy: The energy released when oppositely charged ions come together to form an ionic compound. It is given by Coulomb's law:

$E = -\frac{kQ_1Q_2}{r}$

• Bond Dissociation Energy: The energy required to break a bond in a molecule and separate the atoms in the gas phase.

Ionization Energy and Electron Affinity

• Ionization Energy (IE): The energy required to remove an electron from a gaseous atom or ion.

• Electron Affinity (EA): The energy change when an electron is added to a neutral atom in the gas phase.

• Trends: IE generally increases across a period and decreases down a group. EA becomes more negative across a period.

Electronegativity

• Definition: The ability of an atom in a molecule to attract shared electrons to itself.

• Trends: Electronegativity increases across a period and decreases down a group.

• Bond Polarity: The difference in electronegativity between two atoms determines the polarity of the bond.

Naming Compounds

• Anions and Cations: Cations are positively charged ions; anions are negatively charged ions.

• Ionic Compounds: Composed of cations and anions held together by ionic bonds.

• Acids: Compounds that release H+ ions in solution (e.g., HCl, H2SO4).

• Covalent Compounds: Compounds formed by the sharing of electrons between nonmetals.

Lewis Structures

• Formal Charge: The charge assigned to an atom in a molecule, calculated as:

$\text{Formal Charge} = \text{Valence electrons} - \text{Nonbonding electrons} - \frac{1}{2} \times \text{Bonding electrons}$

• Resonance: When more than one valid Lewis structure can be drawn for a molecule.

• Exceptions to the Octet Rule: Some molecules have atoms with fewer or more than eight electrons (e.g., BF3, SF6).

Molecular Geometry (VSEPR Theory)

• VSEPR: Valence Shell Electron Pair Repulsion theory predicts the shape of molecules based on electron pair repulsions.

• Bond Pairs and Lone Pairs: Bond pairs are shared between atoms; lone pairs are nonbonding electrons on an atom.

• Bond Angles: Determined by the number of electron domains around the central atom.

Dipole Moments

• Bond Moments: The dipole moment of a bond depends on the difference in electronegativity and the distance between atoms.

• Molecular Dipole Moments: The vector sum of individual bond moments; determines if a molecule is polar.

Oxidation-Reduction Reactions

• Assigning Oxidation Numbers: Rules are used to assign oxidation states to atoms in compounds.

• Oxidizing Agents: Substances that gain electrons (are reduced).

• Reducing Agents: Substances that lose electrons (are oxidized).

Chapter 4: Quantum Mechanics and Atomic Structure

Historical Development of Quantum Mechanics

• Line Spectra and Bohr Model: The Bohr model explains the discrete energy levels in the hydrogen atom.

• Hydrogen-Like Atoms: Atoms with only one electron (e.g., He+, Li2+).

• Calculation of Energies: The energy of an electron in a hydrogen atom is given by:

$E_n = -\frac{2.18 \times 10^{-18}\ \text{J}}{n^2}$

• Electromagnetic Spectrum: The range of all possible frequencies of electromagnetic radiation.

• de Broglie Wavelength: The wavelength associated with a particle is given by:

$\lambda = \frac{h}{mv}$

• Uncertainty Principle: It is impossible to simultaneously know the exact position and momentum of a particle:

$\Delta x \cdot \Delta p \geq \frac{h}{4\pi}$

Practice Questions and Applications

• Identifying Chemical Formulas: Recognize the correct formula for compounds such as ammonium cyanate.

• Drawing Lewis Structures: Practice drawing structures for molecules like HNO and identifying formal charges.

• Bond Polarity: Compare bonds such as C-O, B-S, N-F, B-N, and Be-C to determine which is least polar.

• Periodic Trends: Understand how ionization energy and electron affinity change across periods and down groups.

• Resonance Structures: Identify and draw resonance forms for molecules like HSO4-.

• Oxidation Numbers: Assign oxidation numbers to elements in compounds such as V2O5 and HSO4-.

• Quantum Calculations: Calculate the wavelength of light given its frequency, and determine the energy of transitions in the hydrogen atom.

• Photoelectric Effect: Find the maximum wavelength of light that can eject an electron from a metal given its work function.

• Comparing Ionization Energies: Predict and explain trends in ionization energies for elements such as phosphorus and sulfur.

Sample Table: Electronegativity Trends (Inferred from Context)

Element	Electronegativity	Trend
Fluorine (F)	4.0	Highest in Periodic Table
Oxygen (O)	3.5	Second Highest
Sodium (Na)	0.9	Low (Group 1)
Chlorine (Cl)	3.0	High (Halogen)

Additional info: Table values are standard Pauling electronegativity values.

Example: Calculating Wavelength from Frequency

• Given frequency $\nu = 5.4 \times 10^{14}\ \text{s}^{-1}$, calculate wavelength $\lambda$ in nm:

$\lambda = \frac{c}{\nu}$ Where $c = 3.00 \times 10^8\ \text{m/s}$

$\lambda = \frac{3.00 \times 10^8}{5.4 \times 10^{14}} = 5.56 \times 10^{-7}\ \text{m} = 556\ \text{nm}$

Example: Assigning Oxidation Numbers

• In V2O5:

Let $x$ be the oxidation number of V. $2x + 5(-2) = 0$ $2x - 10 = 0$ $2x = 10$ $x = +5$

Example: Energy of Electron Transition in Hydrogen Atom

• For $n = 4$ to $n = 2$ transition:

$E = -2.18 \times 10^{-18} \left( \frac{1}{n_f^2} - \frac{1}{n_i^2} \right)$ $E = -2.18 \times 10^{-18} \left( \frac{1}{2^2} - \frac{1}{4^2} \right)$ $E = -2.18 \times 10^{-18} \left( \frac{1}{4} - \frac{1}{16} \right)$ $E = -2.18 \times 10^{-18} \left( \frac{4-1}{16} \right)$ $E = -2.18 \times 10^{-18} \times \frac{3}{16}$ $E = -4.09 \times 10^{-19}\ \text{J}$

Note: For all calculations, ensure units are consistent and use significant figures as appropriate.